1 / 44

Chemistry 101 : Chap. 8

Chemistry 101 : Chap. 8. Basic Concepts of Chemical Bonding. Chemical Bonds, Lewis Symbols and the Octet Rule (2) Ionic Bonding (3) Covalent Bonding (4) Bond Polarity and Electronegativity (5) Drawing Lewis Structures (6) Resonance Structures (7) Exceptions to the Octet Rule

alia
Download Presentation

Chemistry 101 : Chap. 8

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chemistry 101 : Chap. 8 Basic Concepts of Chemical Bonding • Chemical Bonds, Lewis Symbols and the Octet Rule • (2) Ionic Bonding • (3) Covalent Bonding • (4) Bond Polarity and Electronegativity • (5) Drawing Lewis Structures • (6) Resonance Structures • (7) Exceptions to the Octet Rule • (8) Strengths of Covalent Bonds

  2. Chemical Bonds Chemical bond is formed when two atoms or ions are held together by the attractive force between them.  Ionic Bond : a chemical bond formed between cation and anion  Covalent Bond : a chemical bond formed between two nonmetallic atoms by sharing one or more pairs of electrons.  Metallic Bond : a chemical bond formed when valence electrons of metal atom are attracted by the nuclei of surrounding atoms (electrons are free to move throughout the metal)

  3. Lewis Symbols  Lewis electron dot structure (or Lewis symbol) : Symbol of element surrounded by dots representing the valence electrons in the atom Lewis symbol for sulfur : [Ne]3s23p4 S Maximum 2 electrons on each side  This works only for representative elements (main group) Gilbert N. Lewis (1875-1946)

  4. Lewis Symbols Elements Group e- Configuration Lewis Symbol Hydrogen 1A 1s1 H Helium 8A 1s2He Lithium 1A [He]2s1 Li Berylium 2A [He]2s2 Be Boron 3A [He]2s22p1 B Carbon 4A [He]2s22p2 C

  5. Lewis Symbols Elements Group e- Configuration Lewis Symbol Nitrogen 5A [He]2s22p3 N Oxygen 6A [He]2s22p4 O Fluorine 7A [He]2s22p5 F Neon 8A [He]2s22p6 Ne Note : All four sides of the symbol are equivalent O O =

  6. Lewis Symbols  Elements in the same group of periodic table have the same Lewis symbols F Cl Br I Elements in the same group have the same valence electron configurations For halogen atoms : ns2np5

  7. K Cl Cl Octet Rule  Only the valence electrons are involved in chemical bonding.  Octet Rule : When forming chemical bond, atoms tend to gain, loose or share electrons in order to achieve a complete octet of valence electrons (ns2np6)  same electron configuration as noble gas atom + K+ + [Ar] [Ar] electron configuration: Both ions have an octet of electrons !

  8. Ionic bonding • Ionic Bonding :Cations(metals) andanions (non-metal) combine to form ionic bonds NaCl Alternating positive and negative charges

  9. Ionic bonding  NaCl formation : Na(s) + ½ Cl2(g)  NaCl (s) Hof = -490 kJ Metal : small ionization energy Na(g)  Na+(g) + e- IE = 496 kJ Non-metal : large electron affinity Cl(g) + e- Cl-(g) EA = -349 kJ Removing an electron from Na and transferring it to Cl is NOT exothermic ! Then, why NaCl formation is an exothermic process?

  10. Ionic bonding The main driving force to form ionic bonds is the electrostatic interaction between positive and negative ions. charges of ions distance between ions  Strength of ionic bond depends on Eel  the larger Eel, the stronger the bond  the greater the charges, the stronger the bond the smaller the distance between the charges, the stronger the bond

  11. Ionic bonding The stronger the ionic bond the the melting point higher 1261oC 66, 133 2852oC 66, 140 SrF2 +2, -2 1473oC 113, 133 r1 r2

  12. H H F F Covalent bonding  Covalent bond is formed when two atoms share electrons in order to achieve the electron configuration of the nearest noble gas.  satisfy octet rule Each hydrogen has the electron configuration of He + H H Each fluorine has the electron configuration of Ne + F F

  13. F F Covalent bonding  Lewis dot structure for covalent bonds + H H H H H H single covalent bond F F + F F A shared electron pair is drawn as a dash(two bonding electrons) Unshared electrons are drawn as dots (lone-pair electrons)

  14. Covalent bonding  Example : Draw the Lewis dot structures of H2O and NH3

  15. Double bond Single bond Triple bond C + O + O or O C O O C O or + N N N N N N Covalent bonding  Multiple bond F F + F F F F or

  16. Distance between atoms (bond length) decreases Covalent bonding  Single and Multiple bond X X Bond strength increases X X X X

  17. Drawing Lewis Structure  Things to know before you start to draw Lewis structure  Chemical formulas are often written in the order in which the atoms are connected ex) HCN Hydrogen has only two electrons (shared) and always has only one covalent bond  The central atom is usually written first ex) NH3, CCl4, CHCl3, PCl3

  18. Drawing Lewis Structure  Rules for drawing Lewis structure (1) sum the number of valence electrons from all atoms (2)write the symbols for the atoms andconnect them with a single bond (3)complete the "octet rule" for the atoms bonded to central atom (4) place any left over electrons on the central atom (5) If there are not enough electrons to give the central atom 8 electrons, try multiple bonds.

  19. Drawing Lewis Structure Lewis Structure of NH3 (1) Total number of valence electrons = 5 + 3  1 = 8 (2) Connect atoms with a single bond H N H H and count the number electrons used for single bond = 6 (3) Complete the octets on the atoms bonded to the central atom :done (4) Place remaining electrons (8-6=2) on the central atom H N H H (5) All atoms are satisfying octet. No need to consider multiple bonds

  20. Drawing Lewis Structure Lewis Structure of CO (1) Total number of valence electrons = 4 + 6 =10 (2) Connect atoms with a single bond C O and count the number electrons used for single bond = 2 (3) Complete the octets on the atoms bonded to the central atom (6 electrons are used) C O (4) Place remaining electrons (10-2-6 = 2) on the central atom C O (5) Carbon is NOT satisfying octet rule. Need to have multiple bonds C O C O

  21. Drawing Lewis Structure  Example : Determine the Lewis structure of HCN (1) Total number of valence electrons (2) Connect atoms with a single bond and count the number electrons used for single bond (3) Complete the octets on the atoms bonded to the central atom (4) Place remaining electrons on the central atom. (5) Carbon is NOT satisfying octet rule. Need to have multiple bonds

  22. Drawing Lewis Structure  Example : Determined the Lewis structure of CH2O (1) Total number of valence electrons (2) Connect atoms with a single bond and count the number electrons used for single bond (3) Complete the octets on the atoms bonded to the central atom (4) Place remaining electrons on the central atom. (5) Carbon is NOT satisfying octet rule. Need to have multiple bonds

  23. Drawing Lewis Structure  Example : Determined the Lewis structure of H2O2 (1) Total number of valence electrons (2) Connect atoms with a single bond and count the number electrons used for single bond (3) Complete the octets on the atoms bonded to the central atom (4) Place remaining electrons on the central atom. (5) All atoms are satisfying octet. No need to consider multiple bonds What happens if you choose a different geometry in step (2)?

  24. Drawing Lewis Structure Lewis Structure of ClO3- [ion] (1) Total number of valence electrons = 7 + 63 + 1 = 26 (2) Connect atoms with a single bond and count the number electrons used for single bond = 6 O Cl O O (3) Complete the octets on the atoms bonded to the central atom (18 electrons are used) O Cl O O (4) Place remaining electrons (26-18-6 = 2) on the central atom O Cl O O O Cl O O (5) All atoms are satisfying octet. No need to consider multiple bonds

  25. Drawing Lewis Structure  Example : Determined the Lewis structure of ClO2- (1) Total number of valence electrons (2) Connect atoms with a single bond and count the number electrons used for single bond (3) Complete the octets on the atoms bonded to the central atom (4) Place remaining electrons on the central atom (5) All atoms are satisfying octet. No need to consider multiple bonds

  26. Drawing Lewis Structure : Exceptions • Atoms having fewer than 8 valence electrons : Group IIA and IIIA (mostly Be, B). Example = BeCl2 (1) Total number of valence electrons = 2 + 2  7 = 16 (2) Connect atoms with a single bond and count the number electrons used for single bond = 4 Cl Be Cl (3) Complete the octets on the atoms bonded to the central atom (12 electrons are used) Cl Be Cl (4) Place remaining electrons (16 - 4 -12 = 0) on the central atom : None left (5) Be is not satisfying the octet rule, but no electron is available:

  27. Drawing Lewis Structure : Exceptions • Atoms having more than 8 valence electrons : central atom with n 3, which can use d-orbitals for bonding Example = SF4 (1) Total number of valence electrons = 6 + 4  7 = 34 (2) Connect atoms with a single bond and count the number electrons used for single bond = 8 F F S F F F F (3) Complete the octets on the atoms bonded to the central atom (24 electrons are used) S F F (4) Place remaining electrons (34-8-24 = 2) on the central atom F F (5) S is not satisfying the octet rule (10 electrons) S F F

  28. Drawing Lewis Structure : Exceptions • Molecule having an odd number of valence electrons : Example = NO2 (1) Total number of valence electrons = 5 + 6  2 = 17 (2) Connect atoms with a single bond and count the number electrons used for single bond = 4 O N O (3) Complete the octets on the atoms bonded to the central atom (12 electrons are used) O N O (4) Place remaining electrons (17-4-12 = 1) on the central atom O N O (5) Nitrogen has only 5 electrons. Need to have multiple bonds Free radical O N O

  29. Drawing Lewis Structure : Exceptions  Example : Determine the Lewis structure of BF3, BrF5 and OH

  30. Drawing Lewis Structure : Resonance Lewis Structure of SO3 (1) Total number of valence electrons = 6 + 3  6 = 24 O (2) Connect atoms with a single bond and count the number electrons used for single bond = 6 S O O O (3) Complete the octets on the atoms bonded to the central atom (18 electrons are used) S O O (4) Place remaining electrons on the central atom. No more electron is left (24-6-18=0) O O O (5) Sulfur is NOT satisfying octet rule. Need to have multiple bonds S S S O O O O O O All three S-O bonds have the same length Resonance structures

  31. Drawing Lewis Structure : Resonance  Example : Determine the Lewis structure of O3 and HCO2-

  32. Properties of Covalent Bond  Bond length : The distance between two bonded atoms   bond length Bond length depends on the size of two atoms and the number of covalent bond (single, double or triple) between them.

  33. S Properties of Covalent Bond  Example : Predict which member of each set would have the shortest bond length S

  34. C H (g) Properties of Covalent Bond • Bond Enthalpy: Energy required to completely separate two bonded atoms in gas phase. A short bond is usually harder to break. C (g) + 4 H (g) DH = 1660 kJ/mol (g) per C-H bond: C (g) + H (g) D (C-H)= 1660/4 kJ/mol = 415 kJ/mol

  35. Properties of Covalent Bond Bond enthalpy can be used to estimate the enthalpy change of chemical reactions, Hrxn H2(g) + Cl2(g)  2HCl(g) H = ? H1 H2 Hrxn

  36. Properties of Covalent Bond H1 = D(H-H) + D(Cl-Cl) = 436kJ/mol + 243kJ/mol = 679 kJ/mol H2 = 2 [  D(H-Cl) ] = 2  - 431 kJ/mol = - 862 kJ/mol Hrxn = H1 + H2 = 697kJ/mol – 862 kJ/mol = -183 kJ/mol DHorxn = Σn x Dbroken – Σm x Dformed moles of bonds

  37. Properties of Covalent Bond + + Hrxn bonds broken H – H 1 Cl – Cl 1 bonds formed H – Cl 2 Bond Enthalpy (kJ/mol) H – H 436 Cl – Cl 243 H – Cl 431 Hrxn = 1  436 + 1  243 – 2  431 = - 183 kJ/mol

  38. Properties of Covalent Bond Example : Estimate the Hrxn of following reaction CH4 (g) + 2 O2 (g) → CO2 (g) + 2 H2O (g)

  39. Electronegativity  Electronegativity : A measure of the attraction an atom has for the electron in a bond Metals  low electronegativity Nonmetals  high electronegativity electronegativity scale: Fluorine = 4 (most electronegative)  most strongly attracting electron Cesium = 0.7 (least electronegative)  most easily giving up electron Linus Carl Pauling (1901-1994)

  40. Electronegativity Pauling scale of electronegativity Element EN F 4.0 O 3.5 Cl 3.0 N 3.0 C 2.5 H 2.1

  41. Cl Cl Bond Polarity  Nonpolar covalent bond When two atoms of same element are bonded together, there is equal sharing of the electrons in the bond non-polar covalent bond: equal sharing of electrons

  42. H Cl Bond Polarity  Polar covalent bond When two different elements are bonded together, there is unequal sharing of the electrons in the bond +  - The bonding pair of electrons is pulled toward the chlorine atom (partial charge) polar covalent bond: unequal sharing of electrons

  43. H Cl - Na+ Cl Bond Polarity and Electronegativity +  - ionic bond: electrons are not shared polar covalent bond: unequal sharing of electrons EN = 3.0 – 0.9 = 2.1 EN = 3.0 – 2.1 = 0.9 EN < 0.5 non-polar bond 0.5  EN < 2.0 polar bond EN  2.0 ionic bond

  44. Bond Polarity and Electronegativity  Example : For each pair of bonds, predict which bond is more polar and the partial charge on the atoms (a) Cl – Br Br – F (c) C – H C – O (b) O – F S – F (d) H – O Na – O

More Related